Chemical Bonding: JEE Main and Advanced

Introduction to Chemical Bonding

• Chapter "Chemical Bonding and Molecular Structure" is important for JEE Main and Advanced exams.
• Chapter explains fundamental concepts for organic chemistry.
• Key study aspects: essential concepts, areas to emphasize, and methods for easy understanding.

Why Chemical Bonds Form

• Atoms combine to achieve stability through electron transfer or sharing.

Developments in Bonding Theories

• Scientists' understandings and conclusions are explored through various theories.
• The chapter covers various theories related to chemical bonding.

Achieving Stability

• Atoms combine to become more stable through electron transfer or electron sharing.

Lewis Concept

• Lewis and Kossel independently introduced theories on chemical bonding.
• Atoms combine to complete their octet, mainly, to achieve having eight electrons in their outermost shell.
• Atoms strive to have eight electrons in their outermost shell.

Lewis's Atomic Model

• An atom has an inner core consisting of inner core electrons and the nucleus, called the kernel.
• Valence electrons surround the kernel and they occupy the corners of a cube.
• A cube has eight corners, so when all eight corners are filled with electrons, there is a completed octet.

Sodium and Chlorine example

• Sodium has 11 electrons with a configuration of 2, 8, 1, with its outermost electron occupying a corner of the cube.
• A chlorine atom's configuration is 2, 8, 7, with seven electrons in its outermost shell occupying corners of the cube.

Lewis's Octet Rule

• Each atom tries to achieve eight electrons in its outermost shell completing the octet.
• Sodium tends to give away one electron.
• Chlorine tends to accept one electron to have 8 electrons in its outermost shell.

Lewis Symbols

• Lewis symbols represent valence electrons as dots.
• Lithium has one dot, beryllium has two dots, and boron has three dots.
• Lithium's configuration: 2, 1
• Beryllium's configuration: 2, 2
• Nobel gasses have eight dots representing 8 electrons in outermost shell.

Lewis Dot Symbols in the Periodic Table

• Elements and their dot representation:
• Lithium (Li) has one dot.
• Beryllium (Be) has two dots.
• Boron (B) has three dots.
• Carbon (C) four
• Nitrogen (N) five
• Oxygen (O) six
• Fluorine (F) seven
• Neon (Ne) eight.

Kossel's Approach

• Kossel's approach emphasizes the electrostatic attraction between positive and negative charges.
• The chemical bond concept revolves around electrostatic forces of attraction.

Kossel's Periodic Table Insight

• In the periodic table, highly electronegative halogens and electropositive alkali metals are separated by noble gases, this is key for creating bonds.
• Halogens and alkali metals seek to attain the configuration of noble gases, so they are always looking to make bonds.
• Halogens gain one electron, alkali metals lose one electron, leading to electrostatic attraction and bond formation.
• Lewis and Kossel: Atoms try to attain noble gas configurations.

Octet Rule

• Atoms combine to attain a stable octet.
• In CH4, carbon has four bonds each containing two electrons, resulting in eight electrons and satisfying the octet rule.
• Octet rule is followed by countless molecules.

Limitations of the Octet Rule

• Three main limitations:
  • Central atom with incomplete octet.
  • Odd-electron species.
  • Expanded octet.

Incomplete Octet of Central Atom

• Examples include lithium chloride (LiCl), beryllium hydride (BeH2), and boron trichloride (BCl3).
• Lithium chloride, lithium forms one bond but has only two electrons and not eight.
• BCl3: Boron has three electrons, forming three bonds with chlorine, totaling six electrons around boron instead of eight.
• BeH2: Beryllium forms two bonds, resulting in four electrons around it and not eight.

Odd Electron Species

• NO and Nitrogen Dioxide (NO2) are odd-electron species.
• Nitrogen has two bonds which are 4 electrons, plus three more, resulting in a total of seven electrons in NO.
• These species have an odd number of electrons in the central atom.

Expanded Octet Species

• PCl5 and SF6 are examples.
• Phosphorus has five valence electrons, forming five bonds with chlorine, exceeding the octet rule with 10 electrons.
• SF6: Sulfur has six bonds, and 12 electrons are around sulfur which is more than the octet.
• These expand or super octet molecules cannot form according to Lewis.

Sulfur's Octet Exceptions

• H2SO4, Sulfur completes its octet.
• SCl2 is an example of sulfur obeying the octet rule, with two bonds and four additional electrons totaling eight for sulfur.
• Sulfur can form compounds that obey the octet rule.

Noble Gas Compounds

• Noble gases like xenon and krypton form compounds like XeF2, KrF2, and XeO3.
• The octet rule struggles to explain the reactions of noble gases, since they should always be stable.

Molecular Properties

• It is impossible to predict the shape using the octet.
• Molecular stability as well as energy cannot be spoken about, just facts.

Drawing Lewis Structures

• Lewis structures depict molecules obeying the octet rule.
• Water (H2O) is a common example.
• Oxygen, as a central atom, has eight electrons.
• Example molecules that follow Lewis structures: water, carbon dioxide (O=C=O), carbon tetrachloride (CCl4), ethylene (C2H4), and acetylene (C2H2).
• Lewis provides a representation.

How to Calc Formal Charge

• For polyatomic molecules, formal charge is used to find individual charge as opposed to total net charge.
• A formula for calculating formal charge: Formal Charge = V - L - 1/2 S
  • V is total number of valence electrons.
  • L is total number of non-bonding electrons (lone pair).
  • S is total number of shared electrons (bonding electrons).

Formal Charge Example with Ozone

• Ozone (O3) is used as an example.
• Oxygen, as a central atom, is able to satisfy eight electrons.
• Assign numbers to each oxygen atom in ozone (1, 2, 3).

Calculating Ozone Charge Pt 2

• Calculate the charge for oxygen atom number 2:
  • V (valence electrons) = 6.
  • L (lone pair electrons) = 2.
  • S (shared electrons) = 6 (three bonds).
  • 6- 2 - 1/2 * 6 = +1.
• Now, calculate the charge for oxygen atom number 3:
  • V (valence electrons) = 6.
  • L (lone pair electrons) = 6.
  • S (shared electrons) = 2.
  • 6-6 - 1/2 * 2 = -1.
• You can determine formal charge for various molecules.

Different Types of Bonds

• Ionic
• Covalent
• Hydrogen

Ionic bond basics

• Ionic Bond form when a cation and anion attract you get an ionic bond.

Favorable conditions for ionic bonds

• Low ionization enthalpy means an electron is easy to give away.
• A high negative electron gain enthalpy indicates an electron has an eagerness for receiving electrons.
• High Lattice enthalpy.

Explanation with Lattice Enthalpy

• Lattice Enthalpy: the energy released when one mole of a solid ionic compound forms from its constituent gaseous ions or amount of energy required to break one mole of a solid ionic compound into its constituent gaseous ions.
• Think of Columb's law.
• Charge depends of the ion.
• Size depends of the ion.

Final Summary of ionic bond

• Ionic bonds are with electrostatic attraction force,
• Its common to form a bond in the periodic table with two differing ends.
• Both give and take to complete the electron configuration.

Properties of Ionic Compounds

• Ionic compounds typically have a crystalline structure, high melting and boiling points, and are brittle.
• Will Conduct electricity, but only when melted or in water.

Facts About Soluibility

• CsF (Cesium Flouride)
• CsCl (Cesium Cholride)
• CsBr (Cesium Bromide)
• CsI (Cesium lodide)
• As size increases, Lattice enthralpy decreases so sobolility declines. Also factor in Hydration.
• Ionic: Solluble in water due to both these factors.

Details for Covalent Bonds

• These involve sharing, as opposed to giving or taking.

Bond types

• One pair equal singular bond.
• Two pairs equal Doble bond.
• Three pairs equal Triple bond.

Contions for Convalent Bonds

• Needs a desire for both atoms.
• Negative value to electrons.
• High ionaztion enthralphy to not allow bonds to give out easy.
• small atoms to have a small electronegitivitey.
• Remember that no bond if ever fully can be covarnt or ionic.
• Polarity often over powers and sways to one side.
• Ionic in Covalent bonds is also possible.

Fajan's Rule

• This refers to the force and charecter of said polarity.
• Polarity is diectly correlated to convalent.
• Important to understand Poleerzation: The attaction of ectronic douds.

Factors to Consider on Fajan Side

• Size of the Cation: an atoms nuculues, and inner electronic shell.
• If too big too many shells exist, it gets hard for nucilus so cant attract as it should.
• Therefor, the size of the Cation is inversly poportional to polarizaion for the ion.
• Ex: Bery Clorld, Mag Clirld, store clrold with big ions cant polarize as well.

Fajan's Anions

• Its directly polarizal.
• Catyon: a charge that grows makes the ion able to see the electron cloud.
• Ex Lithim Flo, lithim chloride are going to be the same polorization
• Charge for Catoin direct: as charge gos up. Pulls better the electrons so more effective.
• Aions: Anion charge is directly porpotona.. as charge grows... makes biger. good.

Pseudo Inert Gas Configuration

• Cations vs Anions or Noble gas confliraton.
• Ex: NAchol vs cuck, is used to show it is easier to obtain pseudo.
• NA is good. Cuke is bette.
• Cu+ N3S2 N3PLG T3T10 is easers and more readily pulled, with the 18 electrons

Application of Fujon

• Ag2s vs AgO is for testing power.
• With same chare and sizze there will be different factors.

Anion character part

• the greater electrons, the greater Ionic strength.
• The gigger of a polarity, the more force exists.
• .1 electronegtitivty... if 1.0 greater then ionic, 10 less covalent.

Formaul and Hani

• for polorizaytion

Resonance and its Phenomenon

• When 1 singlet cannot exist to support the property. you need multiple character to show the reality.
• The act of switching traits constantly.
• CO. being the common example, cant exist with two sides so it resonants btween both possibilities.

Terminology

• Conocalal firm vs real firm Conial is to exist, is drawnable but cant exist physically like a drawing Real, we cannot exist but it does physically.
• resonance enefy: ammount of enertgy to reist.. btean the conal, and hybrid

VSP Theory

• Valance Shell, its to see which ones have a low repolsation as possibl.
• Vality depends on the numbers
• Multip bond, has to countd singlty, not doulby or triypled.
• Repouolsotn less and less of a factor

Shape the VS theory side

• 2 / 0... linear
• 2/1 is bend or V shaop.
• important to know how to count.. you do this to determine hybridzaiton as Well.
• You can usse both to find a shape
• The less electronegtivy is the one found the least, or smaller count,

Polar Binds

• Is to go after and want. If no defferince its non Polar bond.
• If they want more, more polizartr
• Think to be as if chare as a distatn
• Measured in Dbu

VSP Theory

• As agnel grows lesser you go
• Direct relsonship with consita.

Molar Orbi Theory

• Its to take acotnm of the oter sode.
• With a nuceali. elextrons revivng
• A molecular orib, will come togereh
• If 2 atomi, = You then have 2 orbit. you either will jave bond of anti bond. Bond less , anti orib better Boding: Adiions with fuctiomn Anit bodn: subt. Bond more bettery!

Formula 2

• What is not listed, or with no start is bondng
• The numbers for bonds matter
• If zero for bond, or negativ... there noth there. No exintance.